Chemical bonds hold molecules together and create temporary connections that are essential to life. Types of chemical bonds including covalent, ionic, and hydrogen bonds and London dispersion forces.

Introduction

Living things are made up of atoms, but in most cases, those atoms aren’t just floating around individually. Instead, they’re usually interacting with other atoms (or groups of atoms). For instance, they might be connected by strong bonds and organized into molecules or crystals. Or they might form temporary, weak bonds with other atoms that they bump into or brush up against. Both the strong bonds that hold molecules together and the weaker bonds that create temporary connections are essential to the chemistry of our bodies, and to the existence of life itself.

Why form chemical bonds? The basic answer is that atoms are always trying to reach the most stable (lowest-energy) state that they can. Many atoms become stable when their valence shell is filled with electrons or when they satisfy the octet rule (by having eight valence electrons). If atoms don’t have this arrangement, they’ll “want” to reach it, and can do so by gaining, losing, or sharing electrons. When atoms share electrons, the shared electrons form a bond that keeps the atoms connected to one another. Similarly, if atoms gain or lose electrons, they become ions (charged atoms) and can form bonds with other ions of opposite charge. These interactions are the basis of the strong chemical bonds in molecules and crystals. But even atoms or molecules that don’t carry full charges can become more stable through temporary charge-based connections with other molecules. These interactions are the basis of the weaker bonds that are essential for biological systems.

Covalent bonds

One major way that atoms can complete their valence shells is by sharing electrons to form covalent bonds. These bonds are generally strong and are common in living systems. For instance, covalent bonds are key to the structure of carbon-based organic molecules like our DNA and proteins. Covalent bonds are also found in smaller inorganic molecules, such as \text H_2\text OH​2​​OH, start subscript, 2, end subscript, O,\text {CO}_2CO​2​​C, O, start subscript, 2, end subscript, and \text {O}_2O​2​​O, start subscript, 2, end subscript. One, two, or three pairs of electrons may be shared between atoms, resulting in single, double, or triple bonds, respectively. The more electrons that are shared between two atoms, the stronger their bond will be.

As an example of covalent bonding, let’s look at water. A single water molecule, \text H_2\text OH​2​​OH, start subscript, 2, end subscript, O, consists of two hydrogen atoms bonded to one oxygen atom. Each hydrogen shares one of its electrons with the oxygen, and the oxygen reciprocally shares one of its electrons with each hydrogen. The shared electrons split their time between the valence shells of the hydrogen and oxygen atoms, providing each atom with something resembling a complete valence shell (two electrons for H, eight for O). This makes a water molecule much more stable than its component atoms would have been on their own.

Hydrogen atoms sharing electrons with an oxygen atom to form covalent bonds, creating a water molecule

Polar covalent bonds

There are two types of covalent bonds: polar and nonpolar. In a polar covalent bond, the electrons are unequally shared by the atoms and spend more time close to one atom than the other. Because of the unequal distribution of electrons between the atoms of different elements, slightly positive (δ+) and slightly negative (δ–) charges develop in different parts of the molecule. This situation, in which two equal and opposite charges are separated in space, is called a dipole. For example, in a water molecule, the bond connecting the oxygen to each hydrogen is a polar bond. Oxygen is a more electronegativeatom than hydrogen, meaning that it attracts shared electrons more strongly, so the oxygen of water bears a partial negative charge (has high electron density), while the hydrogens bear partial positive charges (have low electron density).

More generally, the relative electronegativities of the two atoms in a bond – that is, their tendencies to greedily grab shared electrons – will determine whether a covalent bond is polar or nonpolar. Whenever one element is significantly more electronegative than the other, the bond between them will have some polar character, meaning that one end of it will have a slight positive charge and the other a slight negative charge. Oxygen in particular is quite electronegative (electron-greedy), so be on the lookout for polar bonds in biological molecules containing oxygen atoms.

Nonpolar covalent bonds

Nonpolar covalent bonds form between two atoms of the same element, or between atoms of different elements that share electrons fairly equally. For example, molecular oxygen (\text {O}_2O​2​​O, start subscript, 2, end subscript) is nonpolar because the electrons will be equally shared between the two oxygen atoms. Both oxygens are electron-greedy (electronegative), but they are electron-greedy in the same measure.

Another example of a nonpolar covalent bond is found in methane (\text {CH}_4CH​4​​C, H, start subscript, 4, end subscript). Carbon has four electrons in its outermost shell and needs four more to achieve a stable octet. It gets these by sharing electrons with four hydrogen atoms, each of which provides a single electron. Reciprocally, the hydrogen atoms each need one additional electron to fill their outermost shell, which they receive in the form of shared electrons from carbon. Although carbon and hydrogen do not have exactly the same electronegativity, they are quite similar, so carbon-hydrogen bonds are considered nonpolar.

Table showing water and methane as examples of molecules with polar and nonpolar bonds, respectively

Ionic bonds

Rather than sharing electrons, some atoms achieve stability by gaining or losing one or more electrons. When an atom or molecule gains or loses an electron and becomes positively or negatively charged, it is known as an ion^1​1​​start superscript, 1, end superscript..Anions, or negative ions, are formed through gain of electrons, while cations, or positive ions, are formed through loss of electrons.

Sodium and chlorine atoms provide a good example of ion formation. Sodium (Na) only has one electron in its outer electron shell, so it is easier (more energetically favorable) for sodium to donate that one electron than to find seven more electrons to fill the outer shell. Because of this, sodium tends to lose its one electron, forming Na^+​+​​start superscript, plus, end superscript. Chlorine (Cl), on the other hand, has seven electrons in its outer shell. In this case, it is easier for chlorine to gain one electron than to lose seven, so it tends to take on an electron and become Cl^-​−​​start superscript, minus, end superscript. When sodium and chlorine are combined, sodium will donate its one electron to empty its shell, and chlorine will accept that electron to fill its shell. Both ions now satisfy the octet rule and have complete outermost shells. (In the case of Na^+​+​​start superscript, plus, end superscript, the 2n shell has become the outermost shell because the lone electron from the 3n shell is gone.) Because the number of electrons is no longer equal to the number of protons, each atom is now an ion and has a +1 (Na^+​+​​start superscript, plus, end superscript) or –1 (Cl^-​−​​start superscript, minus, end superscript) charge. In general, the loss of an electron by one atom and gain of an electron by another atom must happen simultaneously: in order for a sodium atom to lose an electron, it must be in the presence of a suitable recipient like a chlorine atom.

Sodium transfers one of its valence electrons to chlorine, resulting in formation of a sodium ion (with no electrons in its 3n shell, meaning a full 2n shell) and a chloride ion (with eight electrons in its 3n shell, giving it a stable octet).

Ionic bonds are electrostatic attractions formed between ions with opposite charges. For instance, the positively charged sodium ion and negatively charged chloride ion shown above will be attracted to each other and form an ionic bond. Compounds created by ionic bonding are known as ionic compounds, and they typically show up in nature as crystals. For instance, we often see sodium chloride (NaCl), or table salt, in crystalline form on the dinner table. Crystalline ionic compounds are made up of many cations and anions that interact with their neighbors in a regular, repeating 3D pattern, and there’s no such thing as a single NaCl molecule; the formula NaCl just represents the overall composition of the crystal structure. Ionic bonds can be very strong in an isolated crystal, but they are much weaker in the presence of water, which readily dissolves many ionic solids due to electrostatic (charge-based) interactions between water molecules and ions.

Certain physiologically important ions, including sodium, potassium, and calcium, are known as electrolytes. These ions are important for nerve impulse conduction, muscle contractions and water balance. Many sports drinks and dietary supplements contain electrolytes that replace ions lost from the body in sweat during exercise.

Hydrogen bonds and London dispersion forces

Covalent and ionic bonds are both typically considered strong bonds. However, other kinds of more temporary bonds can also form between atoms or molecules, and these weaker bonds are very important to the structure of biological molecules like DNA and proteins. Two types of weak bonds often seen in biology are hydrogen bonds and London dispersion forces.

In a polar covalent bond containing hydrogen (e.g., an O-H bond in a water molecule), the hydrogen will have a slight positive charge because the bond electrons are pulled more strongly toward the other element. Because of this slight positive charge, the hydrogen will be attracted to any neighboring negative charges, such as the slight negative charge on the more electronegative atom of a nearby polar covalent bond (often an O or N atom). The interaction produced by this attraction is called a hydrogen bond. This type of bond is quite common, and water molecules in particular undergo extensive hydrogen bonding. Although individual hydrogen bonds are weak and easily broken, numerous hydrogen bonds in combination can be very strong. The hydrogen bond is an example of a dipole-dipole interaction, a broader term for weak interactions between parts of molecules that bear partial charges.

Like hydrogen bonds, London dispersion forces are weak attractions between molecules. However, unlike hydrogen bonds, they can occur between atoms or molecules of any kind, not just molecules with polar bonds. Because electrons are in constant, probabilistic motion, there will be some moments when the electrons of an atom or molecule are clustered together, creating a momentary partial negative charge in one portion of the molecule (and a partial positive charge in another region). This is sometimes called aninstantaneous dipole. If a molecule with this kind of charge imbalance is very close to another molecule, it can actually induce a similar charge redistribution in the second molecule, and the transient positive and negative charges of the two molecules will attract each other. London dispersion forces occur between all types of molecules, and they may be the dominant type of interaction when molecules do not contain polar bonds. Dipole-dipole interactions and London dispersion forces are both types of van der Waals forces, a general term for intermolecular interactions that do not involve covalent bonds or ions.

Although we sometimes talk about these interactions as if they formed clear-cut categories, the reality in a cell is often more of a mixed scenario. Electrostatic interactions between ions, water molecules, polar molecules, and even nonpolar molecules (with induced dipoles) are constantly forming and breaking, and molecules of different types can and will interact with each other. For instance, a Na^+​+​​start superscript, plus, end superscript ion might interact with the negatively charged portion of a water molecule, or with any other negatively charged molecule in the area.